Classical buffer compounds are chemical compounds or agents that resist or minimize changes in hydrogen-ion concentrations when a system of interest is placed under chemical stress or physical stress. Chemical stress includes generation or consumption of hydrogen ions, change in ionic strength and change in solvent properties. Physical stress includes changes in temperature and/or pressure. In many circumstances it is useful or necessary to keep the pH constant or within a narrow range. Buffer compounds are often used in homogeneous aqueous solution, but they may also be used in solid or semisolid mixtures, in non-aqueous polar solvents or in creams, ointments or suspensions. Natural buffer compounds are used in biological systems or living cells to maintain or modulate the proper pH conditions. Natural and synthetic buffer compounds are used commercially in water treatment and sanitation, in agents to control corrosion, in fertilizer mixtures for agriculture, in medicinal formulations and personal care products, in foods and beverages, in fermentation and brewing, in paints and coatings, in human or animal drugs and in many other applications.
pH is a measurement of solvated hydrogen-ion activity, which is directly related to hydrogen ion concentration: −log [H+]. The strength of a monoprotic neutral acid (HA) or monoprotic conjugate acid of a neutral base (HB+) are measured by the pKa values for the respective reactions which are defined by:HA+H2OOH3++A− HB++H2OOH3++BpKa=−log Ka where Ka=[OH3+][A−]/[HA] or Ka=[OH3+][B]/[HB+]Rearranged, these equations give:pH=pKa+log([A−]/[HA]) or pH=pKa+log([B]/[HB+])Each bracketed item in the above equations reflects the “activity” of the respective neutral and ionic species in solution. Activity is a defined thermodynamic term and is related to concentration. At low concentration and low ionic-strength, activity approximately equals concentration. Herein, we use concentration-based pKa values measured at 25° C. at ionic-strength (μ) of 0.10, unless otherwise specified. Such pKa values are more applicable to practical conditions than are thermodynamic pKa values where μ≈0.
The pH range and scale has particular meaning only in a specified solvent of interest and is referenced to certain defined pH standards. The solvation properties of the solvent toward the conjugate acids and bases dictate the degree of ionization and the strength of the acids and bases. Furthermore, most polar solvents have self-ionization properties that limit the strength of strong acids and strong bases dissolved in that solvent. This effect, called the “leveling effect”, is readily apparent in water where the self-ionization constant, pKw, is about 13.8 (25° C., μ=0.1). Under these conditions, neutrality, that is when [OH3+]=[OH−], is defined by pH=0.5*pKw=6.9. When water is the solvent, the acidity (pKa) of solvated OH3+ is about −1.7 while that of solvated OH2 is about +15.5. Dissolved in water, bases stronger than hydroxide simply react with the water forming OH− (hydroxide ion). Similarly, acids stronger than OH3+ react with water forming OH3+. Hence, under practical conditions when the acid/base concentration is about 0.10 M and the solvent concentration (H2O) is about 55.5 M, the leveling effect of water limits the useful pH range for buffering to about 1.0-12.8. Furthermore, other practical limitations in water make working with buffers inconvenient or difficult when operating outside the range 2-12.
An important property of a buffer compound is its buffering capacity, that is, how much OH3+ or OH− can be neutralized by the buffer while minimizing the pH change in solution. The equations above show that the buffering capacity is directly related to (a) the concentration of the buffer and (b) the difference (Δ) between the operating pH and the pKa of the buffer. More buffering capacity can be obtained by (a) increasing the concentration of the buffer compound, (b) by finding a suitable buffer compound so that Δ is smaller or (c) by adjusting the operating pH so that Δ is smaller. A user of buffer compounds must estimate how much acid or base may be generated in the system under study and how much pH variation can be tolerated in order to choose proper buffers and concentrations. Except near the extreme ends of useful pH ranges (pH=2 or 12), the following table allows one to estimate the actual buffering capacity (mM) as a fraction of total buffer compound concentration (CB, mM) at different Δ values. Useful buffering range of a buffer compound is the pH interval within which at least 50% of the maximum buffering capacity is to be maintained. As seen below, the buffering range of a typical monoprotic acid is about 0.96 pH units, that is ±0.48 pH units centered at the pKa value.
Capacity (mM)Δ (pH-pKa)0.500*CB  0.000.443*CB±0.100.400*CB±0.180.387*CB±0.200.333*CB±0.300.285*CB±0.400.250*CB±0.480.240*CB±0.500.201*CB±0.600.167*CB±0.700.137*CB±0.800.112*CB±0.900.091*CB±1.00
Many buffer compounds have been developed for a multitude of uses, including in biological systems. Buffers for use in biological systems face a number of requirements, including low toxicity to living biologic organisms, low metal-binding properties, low sensitivity to pH changes caused by changes in temperature, and specific ionic-charge. Many previously known buffer compounds used in biological systems have suffered from a lack of one or more of these characteristics, and thus there has been an on-going, long-felt need for improved buffer compounds with special, predetermined and controllable properties.
In addition, known buffer compounds, particularly those for use in biological systems, have suffered from an inability to facilely cover the entire pH range. Certain portions of the pH range, for example the range from pH 11 to pH 13, have been notoriously bereft of buffers able to maintain pH within this range.
In addition, because buffer compounds prepared for use in various pH ranges were often chemically quite different, there have been compatibility problems arising from the chemical differences that were inherent in the different buffer compounds needed to obtain buffering at different pH ranges. Ideally, a buffer should control pH only and interact minimally with the constituent molecules under study or of interest in the solution. One important property is its overall charge. If the chemical substances of interest are charged, then a buffer compound of the same charge is used to minimize interaction in solution or mixture. These chemical differences introduce another variable that must be considered when selecting an appropriate buffer compound for a particular application. For these additional reasons, there has been an on-going, long-felt need for improved buffer compounds with special, predetermined and controllable properties.